J.N. Bronsted Essay, Research Paper

In the simple formalism proposed independently by Bronsted and Lowry in 1923, an acid was defined as a proton donor and a base was defined as a proton acceptor. In the simple acid-base reaction shown below, H3O+ is termed a Bronsted Acid, and HO- a Bronsted Base. In writing organic reaction mechanisms, the flow of electrons is often shown using “curved arrows” and in the example shown, the arrows are designed to show that an unshared pair of electrons from hydroxide anion moves to abstract a proton from H3O+, with the simultaneous movement of an electron pair from the bonding orbital to form an unshared pair of electrons on oxygen.

Acid-base reactions are, by definition, equilibria, and the ratio of products and reactants from the proton transfer reaction is given by the equilibrium constant according to the equation shown below.

In the reactions shown above, the two-carbon carboxylic acid, acetic acid (more correctly, ethanoic acid) acts as a Bronsted acid and donates a proton to the Bronsted base, water. The products of the reaction are the carboxylate anion (acetate or ethanoate anion) and H3O+. The equilibrium constant, Ka reflects the position of this equilibrium and the larger the value of Ka, the stronger a given acid will be. For the example shown, Ka 10-4.76, that is, the equilibrium greatly favors the undissociated form of the carboxylic acid. This is because the proton, which is the product of the reaction, is a stronger acid than the carboxylic acid (Ka for the proton is 10+1.74) and equilibria tend to adjust themselves to favor the formation of the most stable forms of all participants. Since Ka is typically a very small number for most common organic acids, it is routinely represented as the negative logarithm of Ka, that is as the pKa. For the example shown above, a Ka of 10-4.76 translates to a pKa of 4.76. Therefore, in this convention, the smaller the value of the pKa, the stronger the acid (the pKa of the proton, Ka = 101.74, is -1.74).

Within a series of simple carboxylic acids, such as the three shown above, the acid strength is controlled largely by the stability of the resulting carboxylate anion. Again, equilibria will always tend to adjust themselves to form the most stable participants; therefore, the more stable a given carboxylate anion, the greater its concentration will be in the equilibrium mixture. In the three equilibria shown, replacing hydrogens on acetic acid with highly electronegative fluorine atoms will result in the formation of more stable carboxylate anions. This is a result of the inductive effect of the highly electronegative fluorine on the carbon-fluorine covalent bonds. As shown below, the bonds will be polarized with most of the electron density directed towards the fluorine. This leaves each fluorine partially negative and the central carbon partially positive. In a similar fashion, this partially positive carbon now has enhanced “effective” electronegativity and will polarize the bond between itself and the carboxylate carbon, removing electron density from the carboxylate group. The net effect is to take the unit negative charge of the carboxylate oxygen and disperse it over the entire molecule; in general, dispersing a charge such as this greatly stabilizes the ion in question. For simple acids, such as those shown, the effect is largely additive; three electronegative atoms stabilize more than two, etc., an the pKas will tend to form a series. For the compounds shown, the pKas are 4.76, 2.66 and 0.23, for acetic acid, fluoroacetic acid and trifluoroacetic acid, respectively. Viewed another way, trifluoroacetic acid is 104.53 (or 33,800) times more acidic than is acetic acid.

A more general description of acids and bases was provided by G.N. Lewis, in 1900. In this definition Lewis Acids are those species which can form a new covalent bond by accepting a pair of electrons and Lewis Bases are species that can from a new covalent bond by donating a pair of electrons. In the example shown below, the Lewis base B donates a pair of electrons to the Lewis acid A to form the covalent compound A-B.

In the acid-base reaction between acetic acid and water, described previously, the Lewis base (the electron donor) is water and the Lewis acid (the actual electron acceptor) is the proton. The real value of the Lewis formulation, however, is in the treatment of non-Bronsted acids and bases. For example, diethyl ether reacts with boron trifluoride to form diethyloxonium fluoroborate (shown below). In this reaction, the Lewis acid (the electron acceptor) is BF3 and the Lewis base (the electron donor) is CH3CH2-O-CH2CH3. Many reactions in organic chemistry involve pre-equilibrium steps which include Lewis acid-base chemistry; an understanding of the Lewis concept will greatly aid in the ability to view these reactions logically and to predict likely reaction pathways and products.